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Chapter 2—Atoms,
Molecules, and Ions
Chemistry: Principles and
Fall 2007
The building blocks of matter:
• Atoms: once thought smallest; we
know now that they’re made of
electrons, protons, and neutrons
• Molecules: groups of atoms bonded
together in a specific way (formula)
• Ions: charged atoms or molecules (+
or -)
Atomic theory
• Dalton developed the atomic model
in 1808.
– An element is composed of tiny particles
called atoms
– In a reaction, atoms move from one
substance to another, but no atom
disappears or is changed into another
– Compounds are formed when atoms of 2
or more elements combine
Were atoms the smallest?
• Was the question for the next ≈100
• Thomson and Rutherford confirmed
the existence of subatomic particles.
• In 1897 Thomson carried out an
experiment involving a cathode
tube—this proved the existence of the
electron (1/2000th of an atom, charge
of -1)
The proton:
• In 1911, Rutherford proved the
existence of the proton with his nowfamous gold-foil experiment
• He shot α-particles at the foil…most
passed through, but some were
reflected back at the fluorescent
The gold foil experiment
“It was as though you
had fired a 15-inch shell
at a piece of tissue
paper, and it bounced
back and hit you.”
--Ernst Rutherford
What did it prove?
• Since α-particles were + charged,
there must be some + charged portion
of the atom
• Rutherford reasoned that a nucleus
with a + charge existed at the center
of the atom…the particles that passed
through the foil missed it…the particles
that bounced back had struck it.
The weird conclusion
• Most of the atom (and most of the
universe) is empty space)
• Since Rutherford, we have learned
that the nucleus contains two types of
– + charged protons almost as big as a
hydrogen atom
– Neutral particles called neutrons, bigger
than a proton
• Most of the mass of an atom, then, is in
the neutron, although the volume is
much larger than the neutron
Atomic number
• The thing that makes an element an
element is its atomic number (the
number of protons)
• The numbers on the periodic table
correspond to the number of protons
in the atoms of that element
• In a neutral atom, the protons and
electrons are equal in number
Mass number
• The mass number of an element is
determined by adding up the protons
and neutrons (electrons are so small
they’re negligible)
• Atoms of the same element (same
atomic number) can differ in mass,
because the number of neutrons can
vary. We call these isotopes
How to write it out
•In the example to the left, the mass
number is the top number.
•The atomic number is the bottom
•How do you determine the number
of neutrons from this information?
•What about the number of
•See example 2.1 for clarification
What is radioactivity?
• Some isotopes are stable, and do not
• Some are unstable, and break down
into nuclei of other elements.
• Stable isotopes have a neutron /
proton ratio falling between 1:1 and
What happens if they fall
• Three things are emitted:
– Beta particles similar to electrons
– Alpha particles, which are heavy helium
nuclei with a +2 charge
– Gamma rays, which are the things that
made Spider-man, the Hulk, and many
other superheroes
The Periodic Table
• Horizontal rows are known as periods. There
are six of them
• Vertical columns are known as groups. There
are 18 of them.
Names to know:
• Group 1: Alkali Metals
• Group 2: Alkaline Earth Metals
• Group17: Halogens
• Group 18: Noble Gases
• In general, metals on the left and
middle, nonmetals on the right,
metalloids in between them: look at
the stairway in your book and know it.
The table in one sentence
• The periodic table is an arrangement of
elements in horizontal order of increasing
atomic number and vertical order of
chemical similarity
• The modern periodic table got its start with
Dmitri Mendeleev, who first began grouping
elements in order of similar chemical
properties—and predicted the discovery of
elements that were later found
What makes a molecule
• 2 or more atoms sticking together
• Most often through covalent bonding,
which means they share electrons.
• We represent molecules through their
formula (NH3) or their structure:
What is an ion?
• An ion is an atom with too few or too
many electrons
• Too few makes it a cation (+ charge)
• Too many makes it an anion (- charge)
• Usually, metals lose electrons (+)
• Usually, nonmetals gain electrons (-)
• The important point---protons never
ever go away.
How they join
• When a metal and nonmetal bond, the
metal gives up its electrons to the nonmetal,
but the neutral charge is maintained:
• Na and Cl come together this way:
• Na gives away an electron (Na +)
• Cl picks that electron up (Cl-)
• So what is the total charge on NaCl?
Ionic compounds
• Ionic compounds are formed by an ionic bond
between a metal and nonmetal.
• The metal gives away electrons, and the nonmetal
picks them up (thus making them both ions
• There are very simple rules for how this happens—
column 1 gives up 1 electron; column 2 gives up 2
• Column 17 accepts 1 electron; column 16 accepts
2; column 15 generally accepts 3
• In general, ions are always trying to get to 8
electrons…noble gas structure
Polyatomic ions
• Polyatomic ions are charged units that
act together in reactions.
• The most common are listed on page
41 in table 2.2. You are responsible to
know their formulae, names, and
charges by one week from today.
Naming Compounds—42-44
• The simplest way is by formula
• Cationstake the name of the metal from
which they come (sodium= Na+
• Metals in the middle can oxidize more than
one way—they get a roman numeral listing
the charge.
• Fe2+ = iron (II)
• Polyatomic ions are memorized
• Monatomic anions get the suffix –ide
• Ionic compounds are named as
follows: first metal + second metalsuffix
• Sodium chloride (NaCl)
• Chromium (III) Nitrate –Cr(NO3)3
Molecular compounds
• Nonmetals and nonmetals don’t form
ionic bonds…they have different rules
• We use the Greek prefix system (page
43), and attach the prefix to each
• N2O5 = dinitrogen pentoxide
Some just have names
• Water = H2O
• Ammonia = NH4
• Acetylene = C2H2
• Nitric / Nitrous Oxide = NO / N2O
The last 2 systems
Hydrocarbons—handled in class
Acids—handled in class